1. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 1 Introductory Chemistry: A Foundation FIFTH EDITION by Steven S. Zumdahl University of Illinois

2. 2 CHEMISTRY 100 Dr. Jimmy Hwang Sections 5 and 7 Tu & Th 9:30a.m.-10:45a.m. Textbook: Zumdahl, Introductory Chemistry, 5 th Edition Office Hours: by appointment 619-421-6700 Ext. 3399 (voicemail) Email: jhwang@swccd.edu

3. 3 Liquids and Solids Chapter 13

4. 4 Overview In Chapter 13, our goals are for the students to: 1. Learn some of the important features of water. 2. Learn about interactions among water molecules. 3. Understand and use heat of fusion and heat of vaporization. 4. Learn about dipole-dipole attraction, hydrogen bonding, and London dispersion forces, and understand the effect of these forces on the properties of liquids. 5. Understand the relationship among vaporization, condensation, and vapor pressure. 6. Learn about various types of crystalline solids. 7.Understand the interparticle forces found crystalline solids. 8. Learn about how the bonding in metals affects metallic properties.

5. 5 Introduction This chapter discusses the properties of liquids and solids. You will learn what makes the particles in solids stay together and why some liquids boil at higher temperatures than others. One liquid of particular importance is water. Liquid water is one of the most important parts of our environment, and we routinely use water to dissolve many kinds of substances. As we will see, water has some unusual properties.

6. 6 Ice, the solid form of water, provides recreation for this ice climber. Source: Vandystadt/Allsport Concepts/Getty Images

7. 7 Wind surfers use liquid water for recreation. Source: Robert Y. Ono/ Corbis

8. 8 A skydiver leaps into the earth's atmosphere containing water in the gaseous state (water vapor). Source: Brian Erler/Taxi/Getty Images

9. 9 Water Colorless, odorless and tasteless Density of ice < than density of liquid water –Not the normal trend –For equal masses of ice & water, ice has the larger volume –For equal volumes of ice & water, water has the larger mass –Water expands when it freezes –Ice floats on water –Density of liquid water = 1.00 g/mL –Density of ice = 0.917 g/mL

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11. 11 Water Freezes at 0°C –At 1 atm, solid at 0°C or below –Normal freezing point = normal melting point Boils at 100°C –At 1 atm, liquid up to 100°C, then turns to steam –Normal boiling point –Boiling point increases as atmospheric pressure increases Temperature stays constant during a state change Relatively large amounts of energy needed to melt solid or boil liquid Liquid’s Specific Heat Capacity = 4.18 J/g-°C

12. 12 Heating Curve As heat added to solid, it first raises the temperature of the solid to the melting point Then added heat goes into melting the solid –Temperature stays at the melting point –Heat of Fusion As more heat added it raises the temperature of the liquid to the boiling point Then added heat goes into boiling the liquid –Temperature stays at the boiling point –Heat of Vaporization As more heat added it raises the temperature of the gas

13. 13 Figure 13.2: The heating/cooling curve for water heated or cooled at a constant rate.

14. 14 Structures of the States of Matter In solids, the molecules have no translational freedom, they are held in place by strong intermolecular attractive forces –may only vibrate In liquids, the molecules have some translational freedom, but not enough to escape their attraction for neighboring molecules –they can slide past one another, rotate as well as vibrate In gases, the molecules have “complete” freedom from each other

15. 15 Figure 13.1: Representations of the gas, liquid, and solid states.

16. 16 Energy Requirements for State Changes In order to change a liquid to a gas, must supply the energy required to overcome the all the intermolecular attractions –Not break bonds (intramolecular forces) The energy required to boil 1 mole of a liquid is called the Heat of Vaporization –  H vaporization = 40.6 kJ/mol for water at 100°C

17. 17 Figure 13.3: Both liquid water and gaseous water contain H 2 O molecules.

18. 18 Figure 13.4: Intramolecular (bonding) forces exist between the atoms in a molecule and hold the molecule together. Intermolecular forces exist between molecules.

19. 19 Energy Requirements for State Changes In order to change a solid to a liquid must supply the energy required to overcome the some of the intermolecular attractions The energy required to melt 1 mole of a solid is called the Heat of Fusion –  H fusion = 6.02 kJ/mol for ice at 0°C

20. 20 Example 13.1 Calculating Energy Changes: Solid to Liquid Calculate the energy required to melt 8.5 g of ice at 0 ºC. The molar hear of fusion for ice is 6.02 kJ/mol. Answer: First change 8.5 g of ice to moles of ice. Next calculate the energy required to melt 0.47 mol H 2 O.

21. 21 Example 13.2 Calculating Energy Changes: Liquid to Gas Calculate the energy (in kJ) required to heat 25 g of liquid water from 25 ºC to 100 C and change it to steam at 100 C. The specific heat capacity of liquid water is 4.18 J/g C, and the molar heat of vaporization of water is 40.6 kJ/mol. Answer: Step 1: Heating to Boiling. Q = s x m x  T where Q = energy required, s = specific heat capacity, m = mass of water, and  T = temperature change

22. 22 Example 13.2 Calculating Energy Changes: Liquid to Gas Step 1. The energy required to heat 25 g of liquid water from 25 ºC to 100 ºC (  T = 75 ºC) is given by Step 2. Vaporization. First find the number of moles of 25 g of water.

23. 23 Example 13.2 Calculating Energy Changes: Liquid to Gas Next calculate the energy required to vaporize 1.4 mol H 2 O given that the molar heat of vaporization of water is 40.6 kJ/mol. The total energy is the sum of the two steps. 7.8 kJ + 57 kJ = 65 kJ Heat from 25 ºC to 100 ºC Change to Vapor

24. 24 Why do Molecules Attract Each Other? Intermolecular attractions are due to attractive forces between opposite charges + Ion to - ion + End of polar molecule to - end of polar molecule –H-bonding especially strong Larger the charge = Stronger attraction Even non-polar molecules have attractions due to opposite charges –London Dispersion Forces

25. 25 Intermolecular Attraction Much weaker than ionic or covalent bonds –Become weaker as distance between molecules increases Dipole-to-Dipole intermolecular attraction due to molecular polarity Hydrogen Bond is a special type of very strong dipole-to-dipole intermolecular attraction –Water has very strong H-bonds London Dispersion Forces are intermolecular attractions between non-polar molecules

26. 26 Dipole-to-Dipole Attractions Size of permanent dipole depends on the bonded atoms and shapes of molecules Individually, dipole-to-dipole attraction is stronger than induced dipole-to-induced dipole attraction –But for larger molecules the London Dispersion forces become more important for predicting the physical properties

27. 27 Figure 13.5: (a) The interaction of two polar molecules. (b) The interaction of many dipoles in a liquid.

28. 28 Hydrogen Bonding Molecules that have -OH or -NH groups have particularly large intermolecular attractions –Also the HF molecule –unusually high melting and boiling points –unusually high solubility in water As electrons are pulled away from H by an electronegative atom, what is left is an unshielded proton that will strongly attract neighboring electrons

29. 29 Figure 13.6: (a) The polar water molecule. (b) Hydrogen bonding among water molecules.

30. 30 Figure 13.7: The boiling points of the covalent hydrides of elements in Group 6.

31. 31 London Dispersion Forces Also Known As Induced Dipoles Caused by electrons on one molecule distorting the electron cloud on another All molecules have them Temporary Size of the London Dispersion Force depends on the number of electrons and shapes of molecules –the larger the molar mass, the larger the induced dipole + -- - - -- -- - - + -- - - -- -- - - + + + + - - - - - - - - - - - - - - - + + + + - - - - - - - - - - - - - - -

32. 32 Figure 13.8(a): Two atoms with spherical electron probability.

33. 33 Figure 13.8(b): The atom on the left develops an instantaneous dipole when more electrons happen to congregate on the left than on the right.

34. 34 Figure 13.8(c): Nonpolar molecules also interact by developing instantaneous dipoles.

35. 35 Attractive Forces and Properties Larger attractive forces between molecules in pure substance means –higher boiling point –higher melting point (though also depends on crystal packing ) Like dissolves Like –Polar molecules dissolve in polar solvents Water, alcohol Molecules with O or N higher solubility in H 2 O due to H- bonding with H 2 O –Non-polar molecules dissolve in non-polar solvents Oils and gasoline

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37. 37 Evaporation Also Known As Vaporization Requires overcoming intermolecular attractions Heat of Vaporization is amount of energy needed to evaporate 1 mole of liquid Condensation the reverse In a closed container, eventually the rate of evaporation and condensation are equal –Equilibrium –In open system, evaporation continues until all liquid evaporated

38. 38 Vapor Pressure Pressure exerted by a vapor in equilibrium with a liquid –Or solid Increases with temperature Larger intermolecular forces = Lower Vapor Pressure Liquid boils when its Vapor Pressure = Atmospheric Pressure –Normal boiling point –Raising external pressure raises boiling point, & visa versa Liquid just poured into open container, little vapor Evaporation faster than Condensation Evaporation as fast as Condensation Equilibrium

39. 39 Figure 13.9: Behavior of a liquid in a closed container.

40. 40 Figure 13.10: (a) It is easy to measure the vapor pressure of a liquid by using a simple barometer of the type shown here. (b) The water vapor pushed the mercury level down 24 mm (760 – 736), so the vapor pressure of water is 24 mm Hg at this temperature. (c) Diethyl ether is much more volatile than water and thus shows a higher vapor pressure.

41. 41 Solids amorphous solids –show no definite structure therefore strengths of intermolecular forces vary over the structure glass, plastic, rubber –tend to soften and melt over a temperature range crystalline solids –orderly, repeating, 3-dimensional pattern –pattern = crystal lattice –melt at one specific temperature

42. 42 Figure 13.11: The regular arrangement of sodium and chloride ions in sodium chloride, a crystalline solid.

43. 43 Figure 13.12: Several crystalline solids: (a) quartz, SiO 2 ; (b) rock salt, NaCl; and (c) iron pyrite, FeS 2.

44. 44 Crystalline Solids

45. 45 Figure 13.13: The classes of crystalline solids.

46. 46 Figure 13.14: Examples of three types of crystalline solids. Only part of the structure is shown in each case.

47. 47 Two examples of atomic solids. The structures of diamond and graphite. In each case only a small part of the entire structure is shown. Diamonds, besides being valuable gemstones, are used for industrial cutting tools. Graphite is a rather soft material useful for writing and for lubricating locks.

48. 48 Types of Crystalline Solids Type of solid depends on type of particle that makes it up Properties of crystalline solid depend on the forces of attraction between the particles Ionic solids are made up of cations and anions arranged in a pattern that maximizes the interaction between ions –Strong forces between ions –Attractions are for all surrounding ions

49. 49 Figure 13.15: The packing of Cl - and Na + ions in solid sodium chloride.

50. 50 Types of Crystalline Solids Molecular Solids are made of molecules arranged in a pattern that maximizes the attractive forces between the molecules –Weak intermolecular attractive forces Atomic Solids of Noble Gases behave like molecular solids –London Dispersion Forces

51. 51 Figure 13.16: (Left) Sulfur crystals contain S 8 molecules. (Right) White phosphorus contains P 4 molecules.

52. 52 S 8 molecules, a molecular solid Sulfur is a non-metallic element which exists in different crystal forms known as allotropes. The main allotropes of sulfur are: Rhombic sulfur is the stable form at room temperature. The crystal shape is shown above. When it is heated slowly above 95.5 C, it is converted to monoclinic sulfur. Both of these forms are insoluble in water but soluble in carbon disulfide, CS 2.

53. 53 S 8 molecules, a molecular solid Both monoclinic and rhombic sulfur are molecular crystals made up of S8 molecules, with 8 sulfur atoms joined in a ring. Crystalline forms of sulfur involve Van der Waals bonding. When sulfur is heated above 113 ºC, it melts to form a pale yellow liquid, which becomes darker and more viscous as the temperature is increased. If liquid sulfur is heated to its boiling point of 445 ºC, and poured into cold water, so-called plastic sulfur is formed, consisting mainly of long chains of sulfur atoms, bound together by covalent bonds:

54. 54 S 8 molecules, a molecular solid When sulfur is heated above 113 ºC, it melts to form a pale yellow liquid, which becomes darker and more viscous as the temperature is increased. If liquid sulfur is heated to its boiling point of 445 ºC, and poured into cold water, so-called plastic sulfur is formed, consisting mainly of long chains of sulfur atoms, bound together by covalent bonds:

55. 55 Figure 13.17: A representation of part of the structure of solid phosphorus, a molecular solid that contains P 4 molecules.

56. 56 Types of Crystalline Solids Atomic Solids of the Non-metals and Metalloids in Group 4A generally have a network of atoms that are covalently bonded together Arrangement depends on VSEPR shape –May be 3-dimensional network (diamond), 2- dimensional network (graphite) or 1-dimensional network (asbestos) –Attraction between atoms is directional

57. 57 Types of Crystalline Solids Atomic Solids that are made of metal atoms –metal atoms release their valence electrons –metal cations fixed in a “sea” of mobile electrons –Leads to strong attractions that are non- directional e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- +++++++++ +++++++++ +++++++++

58. 58 Metallic Bonding & Properties Electrical Conductors because of mobile electrons –electrical current is moving electrons Thermal Conductors because of mobile electrons –small, light, mobile electrons can transfer kinetic energy quicker than larger particles Malleable & Ductile because metal ions can slide past each other in crystal –attractive forces are adjustable because electrons are mobile

59. 59 Metallic Bonding & Properties Luster due to the mobile electrons on the surface absorbing the outside light and then emitting it at the same frequency Melting Points are relatively high due to strong attractions between metal ions and electrons –Stronger ion charge = Larger attraction –Smaller ion size = Larger attraction

60. 60 Metallic Solutions Alloys are mixtures of elements that show metallic properties Substitutional Alloys have some host metal atoms replaced by metal atoms of similar size –Brass Interstitial Alloys have small atoms occupying some of the holes in the crystal lattice of the host metal –Steel

61. 61 Figure 13.18: Two types of alloys. (a) Brass is a substitutional alloy in which copper atoms in the host crystal are replaced by similarly sized zinc atoms. (b) Steel is an interstitial alloy in which carbon atoms occupy interstices (holes) among the closely packed iron atoms.

62. 62 Metal Alloys 1. Substitutional Alloy: some metal atoms replaced by others of similar size. brass = Cu/Zn Substances that have a mixture of elements and metallic properties.

63. 63 Metal Alloys (continued) 2.Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbon 3.Both types: Alloy steels contain a mix of substitutional (carbon) and interstitial (Cr, Mo) alloys.

64. 64 Amount of Carbon Affecting the Properties of Steel Mild steels (< 0.2% carbon), used for nails, cables, and chains. Medium steels (0.2-0.6% carbon), used in rails and structural steel beams. High-carbon steels (0.6-1.5% carbon), used in springs, tools, and cutlery.

65. 65 Steel containing other elements in addition to iron and carbon Alloy steels are mixed interstitial (carbon) and substitutional (other metals) alloys. An example is stainless steel, which has chromium and nickel atoms substituted for some of the atoms. The addition of these metals greatly increases he steel’s resistance to corrosion.